Rechargeable Cells and Batteries (VCE SSCE Chemistry): Revision Notes

Rechargeable Cells and Batteries

What are rechargeable cells?

Rechargeable cells, also known as secondary cells or accumulators, are batteries that can have their discharge reactions reversed. Unlike primary cells that are discarded after use, secondary cells can be recharged and used many hundreds of times. This makes them much more sustainable than single-use batteries.

Common types of rechargeable cells include:

• Lithium-ion cells
• Nickel-metal hydride (NiMH) cells
• Lead-acid batteries

These batteries are found in many everyday devices:

• Mobile phones and laptop computers
• Cameras and portable power tools
• Electric vehicles
• Solar-power energy-storage systems

How the recharging process works

Connecting to a charger

To recharge a cell, you connect it to a charger—a source of electrical energy with a potential difference slightly greater than the cell's voltage. The connection must be made correctly:

• Positive terminal of charger → positive electrode of cell
• Negative terminal of charger → negative electrode of cell

This setup forces the cell's chemical reaction to run in reverse, converting the reaction products back into the original reactants.

Requirements for reversibility

For a cell to be rechargeable, the chemical changes at both electrodes during discharge must be efficiently reversed. Reactions are more likely to be reversible when:

• The electrodes are not damaged during discharge
• The products remain in contact with the electrodes

Energy transformations during discharge and recharge

The key to understanding rechargeable batteries is recognizing that they operate as two different types of electrochemical cells:

During discharge (cell in use):

• Acts as a galvanic cell
• Spontaneous reaction
• Converts chemical energy → electrical energy
• Negative terminal is the anode (oxidation occurs here)
• Positive terminal is the cathode (reduction occurs here)

During recharge (cell being charged):

• Acts as an electrolytic cell
• Non-spontaneous reaction (requires input of energy)
• Converts electrical energy → chemical energy
• Positive terminal becomes the anode (oxidation occurs here)
• Negative terminal becomes the cathode (reduction occurs here)

Lead-acid batteries

History and applications

Lead-acid batteries are the oldest type of rechargeable battery, invented in 1859. They remain the most widely used secondary cell today because they are:

• Relatively cheap
• Reliable
• Capable of providing high currents
• Long-lasting

Most people know lead-acid batteries as car batteries. They serve two main purposes in vehicles:

1. Starting the car's engine
2. Operating electrical accessories when the engine is not running

Once the engine starts, an alternator (powered by the engine) generates electrical energy to run the car's electrical system and recharge the battery.

Structure of a lead-acid battery

A modern lead-acid battery contains six separate cells connected in series. Each cell produces just over 2 V, giving a total potential difference of approximately 12 V.

Components of each cell:

• Positive electrodes: Lead plates packed with lead(IV) oxide ($\text{PbO}_2$)
• Negative electrodes: Lead plates packed with powdered lead (Pb)
• Electrolyte: Sulfuric acid solution (approximately 4 M concentration)
• Separator: Porous material preventing electrode contact

Chemistry during discharge

When a lead-acid battery discharges, it operates as a galvanic cell.

Anode reaction (negative terminal, oxidation):

$\text{Pb(s)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{PbSO}_4\text{(s)} + 2\text{e}^-$

Lead metal is oxidised to lead(II) ions, which immediately combine with sulfate ions to form lead(II) sulfate solid on the electrode surface.

Cathode reaction (positive terminal, reduction):

$\text{PbO}_2\text{(s)} + \text{SO}_4^{2-}\text{(aq)} + 4\text{H}^+\text{(aq)} + 2\text{e}^- \rightarrow \text{PbSO}_4\text{(s)} + 2\text{H}_2\text{O(l)}$

Lead(IV) in lead(IV) oxide is reduced to lead(II), also forming lead(II) sulfate on the electrode.

Overall discharge reaction:

$\text{Pb(s)} + \text{PbO}_2\text{(s)} + 2\text{SO}_4^{2-}\text{(aq)} + 4\text{H}^+\text{(aq)} \rightarrow 2\text{PbSO}_4\text{(s)} + 2\text{H}_2\text{O(l)}$

What happens to the electrolyte?

• Hydrogen ions ($\text{H}^+$) are consumed during discharge
• The pH increases (solution becomes less acidic)
• Sulfate ions are consumed, reducing electrolyte concentration

Chemistry during recharge

When connected to a charger with voltage greater than 12 V, the reactions reverse. The battery now operates as an electrolytic cell.

Anode reaction (positive terminal, oxidation):

$\text{PbSO}_4\text{(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{PbO}_2\text{(s)} + \text{SO}_4^{2-}\text{(aq)} + 4\text{H}^+\text{(aq)} + 2\text{e}^-$

Lead(II) sulfate is oxidised back to lead(IV) oxide.

Cathode reaction (negative terminal, reduction):

$\text{PbSO}_4\text{(s)} + 2\text{e}^- \rightarrow \text{Pb(s)} + \text{SO}_4^{2-}\text{(aq)}$

Lead(II) sulfate is reduced back to lead metal.

Overall recharge reaction:

$2\text{PbSO}_4\text{(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Pb(s)} + \text{PbO}_2\text{(s)} + 2\text{SO}_4^{2-}\text{(aq)} + 4\text{H}^+\text{(aq)}$

The original electrodes and electrolyte composition are reformed, ready for discharge again.

Nickel-based rechargeable batteries

Nickel-cadmium (Ni-Cad) batteries

Towards the end of the 20th century, nickel-cadmium cells were marketed as rechargeable alternatives to common AA and D cells used in toys and small appliances.

Discharge reactions:

Anode: $\text{Cd(s)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Cd(OH)}_2\text{(s)} + 2\text{e}^-$

Cathode: $\text{NiO(OH)(s)} + \text{H}_2\text{O(l)} + \text{e}^- \rightarrow \text{Ni(OH)}_2\text{(s)} + \text{OH}^-\text{(aq)}$

Limitation—the memory effect:

A major problem with Ni-Cad cells is their "memory" issue. If the cells are not fully discharged or fully recharged, their voltage and capacity decrease. This means users had to fully drain and fully charge the batteries to maintain performance.

Nickel-metal hydride (NiMH) batteries

NiMH batteries have largely replaced Ni-Cad batteries because they:

• Have smaller memory effects (less sensitive to partial charging)
• Avoid the health and environmental problems associated with cadmium disposal

In NiMH batteries, cadmium is replaced by a hydrogen-absorbing metal, which stores hydrogen in its structure during operation.

Lithium-ion batteries

Development and applications

Low-density, rechargeable lithium-based batteries with high energy capacity were first sold in 1991. They weren't designed to replace larger NiMH batteries, but rather to enable the miniaturisation of electronic devices.

Lithium-ion batteries enabled the development of:

• Compact laptops and mobile phones
• Digital cameras
• Portable power tools (drills, vacuum cleaners)
• Electric vehicles
• Large-scale energy storage systems

Structure and chemistry

Although different types of lithium-ion batteries exist, they share common features:

• Cathode: Lithium-metal oxide (e.g., lithium cobalt oxide, $\text{LiCoO}_2$)
• Anode: Lithium-carbon compound (e.g., lithium graphite, $\text{LiC}_6$)
• Electrolyte: Allows lithium ions ($\text{Li}^+$) to move through the battery
• Separator membrane: Separates the electrodes while allowing ion movement

Discharge process (battery in use)

When you use a device powered by a lithium-ion battery, the battery operates as a galvanic cell:

Anode reaction (negative terminal, oxidation):

$\text{LiC}_6 \rightarrow \text{Li}^+ + \text{C}_6 + \text{e}^-$

Lithium graphite compound is oxidised, releasing lithium ions that diffuse through the separator to the cathode.

Cathode reaction (positive terminal, reduction):

$\text{CoO}_2 + \text{Li}^+ + \text{e}^- \rightarrow \text{LiCoO}_2$

Cobalt oxide and lithium ions undergo reduction, forming lithium cobalt oxide.

Overall discharge reaction:

$\text{LiC}_6 + \text{CoO}_2 \rightarrow \text{LiCoO}_2 + \text{C}_6$

Charging process

When you purchase a new device with a lithium-ion battery, the battery is only partially charged. You must plug it into a charger and allow full charging before first use.

During charging, the battery operates as an electrolytic cell:

Anode reaction (positive terminal, oxidation):

$\text{LiCoO}_2 \rightarrow \text{CoO}_2 + \text{Li}^+ + \text{e}^-$

Lithium cobalt oxide is oxidised, releasing the lithium ions it absorbed during discharge.

Cathode reaction (negative terminal, reduction):

$\text{Li}^+ + \text{C}_6 + \text{e}^- \rightarrow \text{LiC}_6$

Lithium ions and graphite are reduced to reform the lithium graphite compound.

Overall charging reaction:

$\text{LiCoO}_2 + \text{C}_6 \rightarrow \text{LiC}_6 + \text{CoO}_2$

This process generates the $\text{LiC}_6$ that will react during the next discharge cycle.

Advantages of lithium-ion batteries

Lithium-ion batteries offer several key advantages:

• High voltage output
• High energy storage capacity relative to mass
• Low density (lightweight)
• Can be made very small
• No significant memory effect

Market statistics and future trends

The lithium-ion battery market is experiencing massive growth:

Statistic	Detail
$100 billion	Value of the lithium-ion market in 2022
140 million	Number of electric vehicles predicted worldwide by 2030
11 million tonnes	Mass of lithium batteries likely to reach end-of-life by 2030
5%	Current recycling rate for lithium batteries (only metal components are actually recycled)
100%	Recycling rate for lead from conventional batteries

Sustainability challenges with lithium-ion batteries

While lithium-ion batteries offer many advantages, they raise significant sustainability concerns:

1. Sourcing of lithium

Meeting the demand for lithium is challenging. Australia is the world's leading producer, producing 68,000 tonnes in 2022. However, expanding production to meet the needs of the emerging electric vehicle market will be difficult.

Most Australian lithium comes from Western Australia, with the Talisman mine at Greenbushes being the largest lithium mine in the world.

2. Sourcing of scarce metals

Lithium-ion batteries require several scarce metals:

• Lithium (for both electrodes)
• Cobalt (mainly found in the Democratic Republic of the Congo)
• Nickel

3. Difficulty of recycling

Recycling lithium-ion batteries is much more challenging than recycling lead-acid batteries:

Lead-acid batteries: The electrodes are large pieces of pure lead, relatively easy to separate and recycle. Lead recycling rate is 100%.

Lithium-ion batteries: The electrodes are sophisticated, composite materials. It is not currently possible to isolate individual metals from recycled batteries. Only about 5% of lithium batteries are recycled, and even then, only the metal components are recovered.

4. Safety concerns

Several safety issues have been identified:

Button batteries: Toddlers swallowing button batteries can suffer serious, sometimes fatal, damage.

Fire hazards: In 2021, a fire at the Geelong electricity storage complex took several days to extinguish and emitted toxic fumes. This highlighted the fire risk associated with large-scale lithium battery installations.

Thermal runaway: Damaged lithium-ion batteries can overheat and catch fire, particularly dangerous in confined spaces like vehicles or buildings.

5. Disposal and toxicity

When batteries are added to landfill, the metals they contain (such as cobalt, nickel, and lithium) eventually leach into the surrounding soil and water system. This poses environmental and health risks.

Remember!